History of the Atom
Orginal idea Ancient Greece (400 B.C)
Democritus and Leucippus-Greek philosphers
Aristole-famous philospher
4 elemenets= fire, water, air, earth John Dalton-Teacher ( late 1700)
summarized results of his experiments and those of others Dalton's Atomic Theory
1.) All matter is made of tiny indivisible particles called atoms
2.) Atoms of the same element are identical, those of diffrent elements are diffents
3.)Atoms of diffrent combine in whole number ratios to form ccompounds
4.) Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed Parts of the Atom
*J.J. Thompson-English Physicist (1897)
made a piece of equipment called a carthode ray tube. It is a vacuum tube which means all the air has been pumped out
*Proton- postively charged pieces that are 1840 times heavier than the electron- by E. Goldstein
*Neutron- no charge but the same mass as a proton- by J. Chadwick Rutherfords Model of the Atom
Electromagnetic Radiation
microwave O
infra red R
ultra violet
radio waves Y
x-rays I
visible light
gamma rays
Wavelength
Wavelength is large - energy is small
Wavelength is small - energy is large
Frequency --> how many times/a period of time
Spectroscopes: instruments used to measure the behaviour of radiation Bohr Model of Atom
Ground State Electrons: electrons in the lowest energy level
Drawback of Rutherford Model of Atoms
1) The atom is not stable due to continuous emission of radiation(EMR)
2) Cannot explain the line spectra of Hydrogen like atoms
Quantised: Fixed Energy. Every orbit(shell or energy level) has a fixed value for its energy. (It will not change).
Energy will be absorbed or emitted in discrete quantities.
RAM: Relative Atomic Mass
Unified Atomic Mass: a unit of mass for atoms; 1/12 of the mass of carbon- 12 atom( theortical definition)14= 1/12 of C-12 ( Isotope of C)
Quantum Mechanical Model of an Atom Orbital: It is a 3 dimensional space around the nucleus where there is a very high probability of locating the electrons.
Hesembergs Uncertanity Principle: It is impossible to determine both the position and momentum of an electron with 100% accuracy.
Quantum Numbers: These are a set of 4 integers necessary to specify the energy, position, shape, size and orientation of orbital's to which electrons belong.
Principle Quantum Number "n": It specifies the location and energy of an electron, It is a measure of the effective volume or size of the electron cloud denoted by "n" and can have values 1,2,3,4..........
Azimuthal Quantum Number"l": It determines the shape of the orbital. It takes integral values from 0 (n-1) where "n" id the principle quantum number. When "l" = 0,1,2, the shapes of the orbital's are "s" Spherically symmetrical, "p" dumbbell, and "d" double dumbbell
l
0
1
2
3
4
Orbital
s
p
d
f
g
Shape
Sphercally Symmetrical
Dumbell
Double Dumbell
Magnetic Moment Quantum Number 'm': (2l+1) values It gives the orientation of the orbital's in space. It takes integral values ranging from -l......0.......+l or 's' orbital has got one orientation, 'p' orbital has got 3 orientations, 'd' orbital has got '5' orientatios and 'f' orbital has got 7 orientations.
Spin Quantum Number 's': it indicates the direction in which the electron spins on its own axis or the magnetic property that is associated with it. This value is quantised the two values htat are permitted are +1/2 and -1/2 where g=gyro magnetic constant.
ORBIT
ORBITAL
a well defined circular path around the nucleus in which the electron revolves
a region in 3D space around the nucleus where there is a very high probability of locating the electron
circular in shape
's', 'p' and 'd' orbital's are spherical, dumbell and double dumbell in shape respectively
represents the movement of an electron around the nucleus on one plane
represents the movement of an electron around the nucleus in a 3D space
states with certainty the momentum and position of an electron. Violates the Heisenberg's uncertainty principle
a consequence of Heisenberg's uncertainty priciple. i.e. The posision and momentum cannot be known with certainty
the maximum number of electrons in an orbital is 2n2
the maximum number of electrons in an orbital is 2
Degenerate Orbital's: The orbital's of the same sub shell having equal energy are called degenerate orbital's. e.g. Px, Py, Pz orbital's. The 5 'd' orbital's are degerate but can lose their degeneracy under the infulence of and external electrical or magnetic field.
Nodal Plane: The plane passing through the nucleus on which the probability of finding the electron is almost zero.
Type of Information
Principal Quantum Number 'n'
Azimuthal Quantum Number 'l'
Magnetic Moment Quantum Number 'm'
Spin Quantum Number 's'
Why is it required?
To explain the main lines of spectra
To explain the fine structure of line spectra
To explain the splitting of lines in the magnetic field
Electron Configuration of atoms: rules governing the filling up of electrons in the different orbital's of an atom
Orbital's: arranged in the increasing order of energy which is based on Aufbau order and can be verified using the (n+1) rule.
Aufbau Order: the number of electrons that can be accomodated in an 'orbit' is 2n2 where n is the principal quantum number.
1s,2s,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,6p,7s or 1s2,2s2,3s2,3p6,4s2,3d10,4p6,5s2,4d10,5p6,6s2,4f14,5d10,6p6,7s2
PEP or Pauli's Exclusion Principle: No two electrons in an atom can have all the four quantum numbers alike or an orbital can accommodate a maximum of 2 electrons in an atom.
Hund's Rule of Maximum Multiplicity: Electrons pairing will not take place in orbital's of the same energy (same sub shell) until each orbital is singly filled. *Exceptional Electronic Configuration*Cu, Cr, Mo, Ag, W and Au have got exceptional electronic configuration.
Half filled 'd' orbital configurations are preferred because
there is a extra stability because of the symmetrical distribution of the electrons in the orbitals.
the exchange energy is high for Cr, Mo and W
exchange energy can be calculated using the following equation. where 'n' is the number of electons having parallel spins. In the equation given below 'n' stands for the number of unpaired electrons in the 'd' orbital.
SATP: Standard Ambient Temperature and Pressure. Corresponds to 25 c
STP: Standard Temperature and Pressure used to measure gas density and volume.
Isoelectronic: refers to a group of atoms ot ions having the same number of electrons. For excample, F-, Ne, Na+, are all isoelectronic.
Ionisation Energy:
The energy required to remove an elctron fron the neutral gaseous atom.
It is the energy released when an electron's added to a neutral gaseous atom(usually non-metals)
F+ e-------->F-+ EA
Moving from left to right, the elesctrons are added to the same energy level and there is an increase in nuclear charge. Therefore the size of the atom decreases.
Electronegativity:
It is defined as the tendency of an atom to attract a shared pair of electrons.
Forces of Attraction
INTRAmolecular forces of attraction (inside the molecule)
1.) Covalent Bonds
2.) Ionic Bonds
3.) Co-ordinate Covalent
4.) Metallic Bond
H2 H-H
Cl2 Cl- Cl
H20 H-O-H
INTERmolecular forces of attraction (between the molecule)
History of the Atom
Orginal idea Ancient Greece (400 B.C)
Democritus and Leucippus-Greek philosphers
Aristole-famous philospher
4 elemenets= fire, water, air, earth
John Dalton-Teacher ( late 1700)
summarized results of his experiments and those of others
Dalton's Atomic Theory
1.) All matter is made of tiny indivisible particles called atoms
2.) Atoms of the same element are identical, those of diffrent elements are diffents
3.)Atoms of diffrent combine in whole number ratios to form ccompounds
4.) Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed
Parts of the Atom
*J.J. Thompson-English Physicist (1897)
made a piece of equipment called a carthode ray tube. It is a vacuum tube which means all the air has been pumped out
*Proton- postively charged pieces that are 1840 times heavier than the electron- by E. Goldstein
*Neutron- no charge but the same mass as a proton- by J. Chadwick
Rutherfords Model of the Atom
Electromagnetic Radiation
Wavelength
Wavelength is large - energy is small
Wavelength is small - energy is large
Frequency --> how many times/a period of time
Spectroscopes: instruments used to measure the behaviour of radiation
Bohr Model of Atom
Ground State Electrons: electrons in the lowest energy level
Drawback of Rutherford Model of Atoms
1) The atom is not stable due to continuous emission of radiation(EMR)
2) Cannot explain the line spectra of Hydrogen like atoms
Quantised: Fixed Energy. Every orbit(shell or energy level) has a fixed value for its energy. (It will not change).
Energy will be absorbed or emitted in discrete quantities.
RAM: Relative Atomic Mass
Unified Atomic Mass: a unit of mass for atoms; 1/12 of the mass of carbon- 12 atom( theortical definition)14= 1/12 of C-12 ( Isotope of C)
Quantum Mechanical Model of an Atom
Orbital: It is a 3 dimensional space around the nucleus where there is a very high probability of locating the electrons.
Hesembergs Uncertanity Principle: It is impossible to determine both the position and momentum of an electron with 100% accuracy.
Quantum Numbers: These are a set of 4 integers necessary to specify the energy, position, shape, size and orientation of orbital's to which electrons belong.
Principle Quantum Number "n": It specifies the location and energy of an electron, It is a measure of the effective volume or size of the electron cloud denoted by "n" and can have values 1,2,3,4..........
Azimuthal Quantum Number"l": It determines the shape of the orbital. It takes integral values from 0 (n-1) where "n" id the principle quantum number. When "l" = 0,1,2, the shapes of the orbital's are "s" Spherically symmetrical, "p" dumbbell, and "d" double dumbbell
Magnetic Moment Quantum Number 'm': (2l+1) values It gives the orientation of the orbital's in space. It takes integral values ranging from -l......0.......+l or 's' orbital has got one orientation, 'p' orbital has got 3 orientations, 'd' orbital has got '5' orientatios and 'f' orbital has got 7 orientations.
Spin Quantum Number 's': it indicates the direction in which the electron spins on its own axis or the magnetic property that is associated with it. This value is quantised the two values htat are permitted are +1/2 and -1/2 where g=gyro magnetic constant.
Degenerate Orbital's: The orbital's of the same sub shell having equal energy are called degenerate orbital's. e.g. Px, Py, Pz orbital's. The 5 'd' orbital's are degerate but can lose their degeneracy under the infulence of and external electrical or magnetic field.
Nodal Plane: The plane passing through the nucleus on which the probability of finding the electron is almost zero.
Electron Configuration of atoms: rules governing the filling up of electrons in the different orbital's of an atom
Orbital's: arranged in the increasing order of energy which is based on Aufbau order and can be verified using the (n+1) rule.
Aufbau Order: the number of electrons that can be accomodated in an 'orbit' is 2n2 where n is the principal quantum number.
1s,2s,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,6p,7s or 1s2,2s2,3s2,3p6,4s2,3d10,4p6,5s2,4d10,5p6,6s2,4f14,5d10,6p6,7s2
PEP or Pauli's Exclusion Principle: No two electrons in an atom can have all the four quantum numbers alike or an orbital can accommodate a maximum of 2 electrons in an atom.
Hund's Rule of Maximum Multiplicity: Electrons pairing will not take place in orbital's of the same energy (same sub shell) until each orbital is singly filled.
*Exceptional Electronic Configuration* Cu, Cr, Mo, Ag, W and Au have got exceptional electronic configuration.
Half filled 'd' orbital configurations are preferred because
SATP: Standard Ambient Temperature and Pressure. Corresponds to 25 c
STP: Standard Temperature and Pressure used to measure gas density and volume.
Isoelectronic: refers to a group of atoms ot ions having the same number of electrons. For excample, F-, Ne, Na+, are all isoelectronic.
Ionisation Energy:
The energy required to remove an elctron fron the neutral gaseous atom.
Na+ IE ----------------> Na+ e-
1s2,2s2,2p6,3s1+IE 1s2 ,2s2,2p6 + e-
(less stable ) ( more stable)
Electron Affinity
It is the energy released when an electron's added to a neutral gaseous atom(usually non-metals)
F+ e-------->F-+ EA
Moving from left to right, the elesctrons are added to the same energy level and there is an increase in nuclear charge. Therefore the size of the atom decreases.
Electronegativity:
It is defined as the tendency of an atom to attract a shared pair of electrons.
Forces of Attraction
INTRAmolecular forces of attraction (inside the molecule)
1.) Covalent Bonds
2.) Ionic Bonds
3.) Co-ordinate Covalent
4.) Metallic Bond
H2 H-H
Cl2 Cl- Cl
H20 H-O-H
INTERmolecular forces of attraction (between the molecule)
1.) Hydrogen Bonds (Strongest Force)
2.) Dipole-Dipole Attraction ( Vanderwaals Force)
3.) London's Dispersion ( Weakest Force)
Valence Electrons- The s and p electrons in the outer energy level. It is hte highest occupied eneergy level.
Core electrons- Those in the energy levels below
Electron Configuration for Cations
Metals lose elctrons to attain noble gas configuration
They make positive ions ( cations)